In this lecture we're going to continue to understand the structure of an atom by beginning our study of how the electrons move in the great space around the nucleus in the atom. You'll recall that in our previous lecture, we discussed the concept of the atomic structure, and in particular the nuclear structure of an atom. The idea here is that there's a tiny nucleus surrounded by vast, open space in which the electrons move. Let's sort of try to draw a picture of what that looks like. So we've got this very tiny, positively charged nucleus here. It turns out that the diameter of this nucleus is o of the order of 10 to the minus 15. To 10 to the minus 14 meters, so that's extremely small. How is that in comparison to the atom? Well, the radius of the atom turns out to be much larger than that. So we have this vast empty space out here in which the electrons move, and the radius of the atom itself turns out to be more of the order of 10 to the minus 10. To 10 to the minus 9 meters. So you'll notice that there's a difference in these ratios that's of the order of between 10,000 and maybe a 100,000. Those are fairly abstract numbers. We certainly don't understand how large 10 to the minus 15, or 10 to the minus 14 meters might be. But what we can say is some metaphor to that, or some analogy to what that might look like. Here's a picture from the concept development study. That might help you have an idea just how tiny the nucleus is, relative to the size of the rest of the atom. So here's a map of the city of Houston, and what we've done here, Rice University is right in the middle of this map about this location here. We've put a basketball down right in the middle of the Rice Campus here, of course this basketball is not really drawn to scale, it should be much, much smaller, but then we couldn't actually see it. So, imagine the size of a basketball, if that is the size of a nucleus, then the size of the atom is about this radius that we've drawn out here, which is about 12 kilometers large. If you're familiar with the city of Huston, This highway here, 610, is what we call the loop, and has a radius of around seven or eight miles, and in comparison then, this circle is about the size of that, in comparison, the size of the nucleus is something on the order of a basketball. So that's actually extremely small. So one of the things we've already learned then is that the nucleus is positively charged, but is quite tiny. But virtually all of the mass of the atom, is inside of that tiny positively charged mass. And the electrons are moving in this huge open space then out here, surrounding that nucleus. In addition, we ran an experiment in the last study where we learned something about the atomic number, the atomic number was simply a number we assigned based upon the mass rank of the elements. But we've now shown that the atomic number is actually equal to the positive charge on the nucleus and since atoms are neutral, then it's also equal to the number of electrons which are outside of the nucleus. Now that we know that information, it seems reasonable that we should assume that the chemistry of a particular atom should be determined. By what these numbers of positive charges are. That it might seem to make sense to us since each atom has its own unique atomic number and has its own unique chemistry that the chemistry of an atom is related to the numbers of positive charges and the number of negative charges. Given that's the case. Let's actually make a comparison now, of some elements that have similar atomic numbers, and remember if they have similar atomic numbers then they have similar numbers of positive charges in their nuclei. And similar numbers of pos-, of, of, of negative charges, for the numbers of electrons moving around. Here are five consecutive elements, that we've picked out here. That have very similar atomic numbers, 8, 9, 10, 11, and 12. And therefore, similar positive charges. What we see in this particular graph is that although they have very similar positive charges, the elements are very different. Oxygen and fluorine are both very reactive diatomic ga, gases. By contrast neon, which has only one more positive charge than fluorine, is very non-reactive, monatomic, or atomic gas. It doesn't really form any compounds at all. And although sodium has only one more positive charge than neon, it's a completely different material from neon. It is a highly reactive, solid metal, and likewise magnesium is also a solid reactive metal, although it's less reactive than sodium is. What we can conclude from this comparison is in fact, knowing that we have similar numbers of electrons, does not tell us that we have similar kinds of elements. To the contrary, actually. Similar numbers of electrons might produce very different chemical properties. But that seems contrary to what we were thinking before that somehow or another, the number of electrons should have something to do with the properties of the particular atom. Let's try then to look for elements that do have similar kinds of properties. Here's a group of them. We're going to take this particular group here of fluorine, chlorine, bromine and iodine. Turns out that they are quite similar. They are all diatomic elements. Let's see. Fluorine and Chlorine are both gases at room temperature. Bromine is a liquid at room temperature, but it boils at a temperature not far above room temperature. Iodine's actually a solid which sublimes into the gas phase at temperatures not far above room temperature. So notice that these things are fairly similar but notice also that they have very different atomic numbers. That means they have very different positive charges and they have very different negative charges. So we can conclude that we can get very similar properties from very different numbers of electrons. That tells us that the structure of the atom arrangement of electrons about the atom must somehow or another be a bit more subtle and complicated than what we have learned in our structure so far. So what were going to do next is to take a look at some of the elements of the periodic table and noticed that infect those elements can often be grouped. According to their similar properties. That actually is what we were illustrating in the previous slide here. This particular set of elements here turns out to be a group. The elements fluorine, chlorine, bromine and iodine, with very similar chemical and physical properties, they form very similar kinds of chemical compounds. Because they are so similar, we give them a name and we call them the halogens. There are other elements such as say sodium, potassium, rubidium, say lithium if I list them in the order of their increasing atomic masses that are all relatively soft highly reactive metals and we refer to these as the alkaline metals. And of course the simplest group perhaps would be helium, neon, argon, krypton and so forth, which we refer to as the noble gases in recognition of the fact that they are either unreactive or extremely unlikely to form compounds. So the first step in our process here is to recognize that elements can be grouped according to similar properties. The next result is the very surprising one, which is that if we look at the elements in the groups, for example this fluorine, chlorine, bromine, or iodine. They appear in the Periodic Table periodically. So let's actually pull up a table here that illustrates that that's the case. This is actually from the concept development study on the shell model of an atom. So let's go look for elements that have similar kinds of properties here. Here, for example, is fluorine. And here, for example, is chlorine. And, let's see, slide this up a little up further so that we can see it, and down below, we'll actually find bromine. And if we notice where these elements are, in each case they, the elements in the group which are called halogens immediately precede the elements which we call the noble gases. Notice fluorine precedes neon, chlorine precedes argon and bromine precedes krypton down below. And after each noble gas, including if we look up, back above to helium. Immediately after helium is a member of the alkaline metals, and immediately after neon is a member of the alkaline metals. And immediately after argon is a member of the alkaline metals, and I'll look down at the bottom, we have not shown this, the next element down below is another alkaline metal, called rubidium. The point here is that the different groups of elements such as for example lithium, and sodium, and potassium appear recurrently, they appear periodically in the periodic table. That actually is an example of something we call the periodic law. The periodic law says. That if we rank the elements by their atomic number the physical and chemical properties of the elements are all periodic functions of that atomic number. That is, the properties recur over and over again as we go through the list order of the elements. That causes to give rise to something called the periodic table of the elements. You can get access to a periodic table in lots of different locations, but it's a beautiful construct that allows us to look at, the relationship of the different elements to each other. This particular one is taken from the Royal Society of Chemistry's website. They call it their visual elements periodic table. Notice in the periodic table here, what we've done is to list the periodic elements moving across the periodic table here as we increase the atomic number. But we've also stacked up the elements which have similar chemical and physical properties. So fluorine, chlorine, bromine, and iodine are the four elements we listed as halogens. Helium, neon, argon, krypton, xenon, are the elements we listed as noble gases. Lithium, sodium, potassium, rubidium are the elements that we listed as the alkaline metals. In each case, each column of the periodic table corresponds to a set of elements which have somewhat similar chemical and physical properties. Whereas the rows are simply increasing numbers of, of atomic numbers, the increasing number in the mass order. Let's go back to our slides here then. And say that we now discovered that although the, it's not a simple relationship between the number of electrons and the number of protons in an atom, it appears to be a periodic function of the number of electrons and protons in an atom which determines its chemical and physical properties. We're going to consider one very important example of the elements and their periodicity which is something called ionization energy. The ionization energy as expressed here it is the amount of energy required to take an electron away from an atom. Remember that our atom looks something like this. Positively charged nucleus surrounded in some large space here in some way or another with electrons flying around. And the idea here is what if I were to try to remove that electron. If I took the electron away what would remain behind would be a positive ion. But it would cost us energy to be able to pull that electron away from the atom and that's a consequence of the fact that there's a force of attraction between the electron and the positively charged nucleus, very much like these two magnets that I've got here. These are magnets you get in the toy store. In this case, if I want to pull these two magnets apart as they are attracted to each other, I have to add energy to them to pull them apart. I have to do some work. Similarly if I let the atoms or the magnets come back together, actually energy is released as they attract each other. So, energy is absorbed when I pull them apart and is released when I bring them back together. Therefore, if I want to remove an electron from an atom I am going to have to apply an ionization energy. That is the minimum amount of energy which is required to remove an electron. Without going into details about how we are going to measure the ionization energy let's actually just look at what the data appear to be as we look at the ionization energies. Here's a graph of the minimum energy required to remove an electron from each of the elements of the periodic table and we've for a set of them and we've simply graphed them as a function now of the atomic number. And notice that you actually see the groups in the periodic table in kind of conspicuous and similar sorts of locations. So, for example, right at the top of each of these, we see the noble gases, and right at the bottom of the chart in each case, we see the alkali metals. So there appears to be groups that have similar kinds of properties so the periodic law is really quite conspicuous here. Well what's going on in this graph in addition? Let's remove these markings here to make them a little easier to see. Let's go back and look at our list of elements. Or actually wanted to look at the periodic table here. And bring it across to where we can compare the periodic table to the ionization energies here. Notice in each case that as we move across the periodic table from, say. Lithium, beryllium, boron, carbon, all the way out to neon. That, that corresponds to this particular stretch moving from lithium to neon. And then if we go to the next row of the periodic table, sodium, magnesium, going all the way across to argon, that corresponds to this rise. And notice we keep getting these periodic rises as we move from the alkali metals all the way across to the noble gases. Within an addition we get this rather precipitous drop each time we move from a noble gas down to the subsequent alkali metal even back over here as well. So there are some clear repeating patterns in this data that have something to do with the arrangement of the electrons in the table. For now we're simply going to draw two conclusions. The first is that when you measure the ionization energies of the atoms, the periodic law is really quite clear in this state. This actually looks like a periodic function. And further more we can easily see the appearance of the groups of elements we described earlier in this lecture appearing in this data. Why the ionization energy looks this way what we're going to take up in the next lecture when we start to understand exactly how this data tells us the arrangement of the electrons about the nucleus.
